Inorganic Chemistry - A Model Chapter

Chapter 6 continued-

As the possibility of a chemical reaction is intimately linked to the nature of chemical bonding, it will be easier to predict the course of the reaction if, only ionic bonding is present than when there is great decrease in ionicity (ie increase in covalency). According to Seel, an ionic reaction will take place only if the reaction leads to the formation of new groupings where the bonding is stronger than the bonding in the reactants. As the products have stronger bonding_.than the reactants, the overall molar volume⁽⁸⁾ of the products will be smaller than the overall molar volumes of the original compounds. Decrease in molar volume may occur if there is a possibility for better matching of the anion and the cation; for appropriate matching of the anion and the cation favours stability.

Consider the following reactions:

      KF            +    LiBr                              KBr            +         LiF                        (1)

     NaBr        +    KCl                               NaCl             +          KBr                        (2)

     KBr          +    AgF                                 KF               +          AgBr                      (3)

     H₂O          +    HI                                   OH₃⁺I^-                                                   (4)

    Na₂SO₄     +    BaCl₂                                         BaSO₄       +       2NaCl                     (5)

    2NaF         +    CaCl₂                          2NaC         +       CaF₂                                  (6)

    K₃(FeCl₆)  +   6KF                             K₃(FeF₆)    +       6KCl                      (7)

    2KCl          +   SnCl₄                              K₂(SnCl₆)                                          (8)

    KF             +   MgF2                              KMgF₃                                                  (9)

In the reaction (1) the electronegativity difference between potassium and fluorine is higher than the electronegativity difference between potassium and bromine. Despite this, the formation of KBr is favoured because K⁺ ion (rK⁺ = 1.33 A^°) is more equal in size to Br^- ion (rBr^. = 1.98A^°) than the Li⁺ion (rLi⁺ = 0.68A^°) to the Br^- ion. In the reaction (2) also, though all the ions will remain free in solution, the forward reaction is favoured because of the proper matching (in size) of the anion and the cation ie the biggest combines with the biggest, the smallest combines with the smallest. Numerous reactions where matching of size contributes to the stability of the product compounds may be cited. For instance, when heated with carbon, both Li and Na react to

⁽⁸⁾ The molar volume is the number of milliliter  occupied by the gram formula weight of a compound ie. the quotient of the gram formula weight and density.

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form acetylides - Li₂C₂ and Na₂C₂. Though other alkali metals react with carbon, only non- stoichiometric interstitial compounds are obtained.

As BF^-₄ ion has a tetrahedral symmetry, it has greater stability than BF₃. Hence, BF₃ generally accelerates the ionization of R-F (where R is an alkyl group) forming the stable BF₄^- ion. But the same compound, BF₃, does not accelerate the ionization of R-Cl because (BF₃C1)^- ion is less stable than BF₃. Symmetry considerations, in fact, show that (BF₃Cl)^- ion must be less stable than BF₃. But the reverse is true with AgF as catalyst because AgCl is more stable than AgF. The stability of BF₄^- ion indicates that ions with a cluster of atoms also possess a special stability when there is proper matching of atoms.

In the equation (3) in spite of the fact that the K⁺ ion (1.33 A^°) is bigger than the Ag⁺ ion (1.13A^°) the forward reaction is strongly favoured because AgBr is a precipitate. It indicates that the cations with higher polarizability tend to combine with the anions of higher polarizability. (Pseudo inert gas configuration is more polarizable than inert gas configuration). Thus, when a solution of silver nitrate is titrated against a solution containing I^-, Br^- and Cl^- ions, I^- ions will be consumed first followed by Br^- and Cl^- ions, the influence of polarizability of the cation being generally greater than that of electronegativity difference and ionic size. (The low solubility of the silver halides and of the sulphides of metals with highly occupied d shells is attributed to dispersion forces also). PbC1₂, unlike BaC1₂, is insoluble in cold water. It is also presumably due to the polarizability of the highly occupied d shell (in the n-1 level) in Pb. PbF₂ and AgF are soluble in water mainly due to the absence of stable d orbitals in the fluorine atom for π overlapping⁽⁹⁾. In the equation (4) it can be seen that the highly polarizable iodine atom fails to keep the very small hydrogen atom with it since the hydrogen atom is not polarizable. Iodide ion, it must be pointed out, is a very effective nucleophile. In fact, it is most easily oxidized even by such feeble oxidizing agents as Cu²⁺.

⁽⁹⁾ Though the solvents with high dielectric constants will be better for the dissolution of the ionic compounds, solubility of a substance in a particular solvent, however, cannot be the sole basis for interpreting the 'nature of the bond in the substance. Thus, BaC1₂, (NH₄)₂ SO₄, NaOH etc may be precipitated from liquid NH₃, CuSO₄ from acetic acid etc. Similarly, although the dielectric constants of HF and H₂O are very similar, the former dissolves for fewer salts presumably because of the inadequate solvation of the cation in liquid HF. However, it can be generalized that the higher the lattice energy of a particular crystal, the less likely is that crystal to dissolve in water.

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Thirdly, as the lattice energy increases with decrease in the size of the anion, decrease in the size of the cation and increase in the charge of the anion, in a system with more than two kinds of ions, the cation having the highest charge density combines with the anion having the highest charge density; and the cation having the lowest charge density combines with the anion having the lowest charge density, the influence of charge density being generally higher than that of electronegativity difference, ionic size and polarizability. This type of combination is favoured because, here, the electrostatic attractive forces between the oppositely charged ions are maxi­mum. The above generalization is especially true when the cation has inert gas configuration. Thus in the reaction (5) the charge density of Ba²⁺ being higher than that of Na⁺, Ba²⁺ combines with the ion of highest opposite charge density ie with the S0₄^2- ion. Reaction (6) also is an example where matching of charge manifests. The charge density of Ca²⁺ ion (rCa²⁺ = 1.0A^°) being higher than that of Na^t ion (rNa⁺ = 0.98A^°), Ca²⁺ ion combines with the ion of highest opposite charge density namely F^- ion. To cite more examples, it can be noticed that in the series of compounds CsI (m.p.621^°C) - NaCl(m.p.801^°C) - LiF (m.p.850^°C) and in the series BaO (m.p.1925^°C) ­CaO (m.p.2600^°C) -MgO (sublimes at 2770^°C) the melting points show a gradual increase as the lattice energies increase. If the solubility of the salt is considered, salts like Na₂SO₄, BaC1₂, Ba(NO₃)₂, LiC1O₄, KF, PbF₂, AgF are soluble in water. These salts have weak lattices since there is no effective matching of charges. In contrast, salts like BaSO₄, LiF, PbCl₂, AgC1 are insoluble in water. In the formation of ionic complexes and in their reactions also, the smallest cation (ie the most highly charged cation) will combine with the smallest anion. This is because of the simple reason that attractive forces are greater when the charge densities of the ions are greater. Thus, in the reaction (7) the chlorine atom is substituted by the fluorine atom.

If, for instance, KCl and SnCl₄ are dissolved together in water, the solution will have initially K⁺, Sn⁴⁺ and Cl^- ions. Sn⁴⁺ ions and Cl^- ions will aggregate to give the complex SnC1₆^2- ion (equation 8). This is simply because attractive forces between the anion (here C1^-) and the small highly charged cation (here Sn⁴⁺) are much stronger than those between anions and larger cations of low charges. When the attractive forces are weak the ion cannot approach so

Note: Anhydrous inorganic compounds do not contain water either adsorbed on their surface or combined as water of crystallization. It should not be confused with anhydride which is a chemical compound derived from an acid by elimination of a molecule of water. Thus, SO₃ is the anhydride of H₂SO₄ .

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closely. Thus, the formation of complex ions such as (KC1₆)^5- cannot be simply envisaged.

In the series BeSO₄ - MgSO₄ - CaSO₄ - SrSO₄ - BaSO₄, since the charge densities of M²⁺ ions decrease, the covalent character of the compounds must decrease and the solubilities of the compounds must increase. But, actually, the solubilities decrease at the same time the melting points increase, indicat­ing the increase in lattice energies. In BaSO₄ there is appropriate matching with regard to charge and size. BeSO₄, MgSO₄ and CaSO₄ are hydrated. The hydrated water molecules, apart from preventing the cation from polarizing the anion, serve the purpose of increasing the size of the cation so that the salt attains appropriate matching with regard to charge density and size (refer chapter 30). Thus, anhydrous CaSO₄ is slightly soluble in water but the dihydrate is a pure precipitate. It loses 1/2 H₂O at 128^°C and becomes anhydrous only at 163^°C. NaF and NaCl are anhydrous. But, NaBr forms dihydrate with the release of 4.6k cals mol^t. Hydration, due to the obvious reason, tends to increase with decrease in the size of cation eg BeSO₄ 4H₂O and BeCl₂ 4H₂O.

Further, as the big cations with the large anions are more stable, complex cations such as hydrates and amines are most stable with large anions. But salts like NaClO₄, NaIO₃, NaC1O₃, NaC1O are not hydrated. It is attributed to the multipole character of oxyanions.

In the reaction (9) the forward reaction is favoured because potassium magnesium fluoride crystallizes in the particularly economical perovskite (calcium titanate) lattice, the influence of lattice being greater than other factors. Here, the molecule has the best compromise of maximizing electron-nucleus attraction through appropriate matching and minimizing electron­-electron repulsion. Thus, this reaction results in the decrease in the molar volume of the reactants, the molar volumes of KF, MgF₂, and KMgF₃ being 22.8, 19.6 and 38.6 respectively. Here, it must be noted that the radius of K⁺ ion (1.33A^°) is nearly equal to the radius of the F^- ion (1.33A^°). On the basis of the above interpretation it can be predicted that the reaction of the type

          ICl  +    F₂                                    IClF₂

may not take place because of the probable increase in the molar volume. Indeed, interhalogen compounds of the type IBrC1₂, IBrF₂, IBrF₄, IC1F₆ do not exist. In general, in a given medium several factors influence the course of the chemical reaction. The order of influence of such factors approximately

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being : electronegativity difference < ionic size < polarizability < charge density and < lattice. ⁽¹⁰⁾

The above generalization gives an insight into the concept of hard and soft

acids and bases. The hard acids such as H⁺, Li⁺, Mg²⁺,                    Ti⁴⁺, BF₃ tend
to preferentially polarize those ligands which have greater electron density or electron donating power. Thus, the tendency of the hard acids to complex with bases is in the order

N             >>           P              >              As           >              Sb

O             >>           S              >              Se            >              Br

F              >              Cl            >              Br            >              I

The soft acids such as Cu⁺, Ag⁺, Hg²⁺, GaI₃, RSe⁺ tend to polarize tho se ligands which have greater polarizability. Thus, the tendency of the soft acids to complex with ligands is in the order

N             <<           P              >              As           >              Sb

O             <<           S              <              Se            ~              Te

F              <              Cl            <              Br            <              I

Phosphorus and sulphur atoms possess not only lone pairs but also very stable d orbitals for TE overlapping. In fact, back co-ordination of the metal greatly strengthens the binding in complexes and it is presumably responsible for the fact that phosphines form much more stable complexes with metals than the corresponding amines, although the latter are stronger bases. Another equally probable reason why phosphines form much more stable complexes with metals than the corresponding amines is the singularly great tendency of phosphorus atom to have tetrahedral configuration (refer chapter 11). However, the general observation here is that ligands with strong donor properties such as water, ammonia favour the stability of higher valence states and ligands with π bond acceptor properties but weak o bond donor properties favour the stability of lower valence states. It appears that the electrostatic attractions are important in the former and the polarizability factors in the latter.

⁽¹⁰⁾ This generalization has not previously been published. Further, it is ac­knowledged that this generalization does not seek to disregard changes in free energy, heat content and entropy which are generally considered as the qualities of impor­tance for an inorganic chemist.

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K. If the atoms and molecules of a compound are arranged in an ap­propriate way with great precision and perfection, certain materials lose their electrical resistance, ie they become superconducting, below a particular temperature called critical temperature.

In order to increase the critical temperature, researchers usually "sandwich" copper oxide (with a little calcium) between layers of strontium and bismuth. But improvement ceases when they put more than three layers of copper oxide in the sandwhich. Many researchers believe that progress is being blocked by the way the material is made, rather than any basic law of physics.

Now * Michel Lagues and his colleagues^-at Paris's Elite College of Industrial physics and chemistry claim that they noticed superconductivity at 250K in a ceramic material made of bismuth, strontium, calcium and copper oxide. This material is not completely stable as the superconductivity dropped to around 200K after two weeks. Lagues's material consists of eight stacked layers of copper oxide, one atom thick, sand-wiched between each layer of bismuth and strontium. He carefully built up a stack of about 100 layers to make a film about 30 nanometres thick and published his work in the journal Science, Dec.17, 1993.

Lagues built up the layers with scrupulous precision. He inspected each one for imperfections and threw out flawed samples. He was able to reduce the number of flaws in his layers using a laser technique that allowed him to work at very low temperatures. This showed the movement of atoms so that they more or less ceased vibrating, cutting the number of imperfections in the thin films.

Lagues's ceramic material is a great leap in the preparation of supercon­ducting materials (refer chaper 21) although the results need to be confirmed.

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Further Reading :

I. Akhmetov. N; Inorganic Chemistry, MIR Publishers, Moscow. 1973.

2.     Clyde Day, M JR andJoelSelbin, Theoretical Inorganic C hemistry, Affiliated East-West Press, Pvt. Ltd, New Delhi. 1977 .

3.     Edwin S. Gould, Inorganic Reactions and Structure, Holt, Rinehart & Winston, Inc, New York.

4.     Esmarch S. Gilreath, Fundamental Concepts of Inorganic Chemistry, Mc­Graw - Hill Book Company Inc. 1958 .

5.     George E. Ryschkewitsch, Chemical Bonding and the Geometry of Molecules, Chapman & Hall Ltd, 11 New Fetter Lane, London EC4. 1965.

6.     Jerry March, Advanced Organic Chemistry, Second Edition, McGraw H'll International Book Company, New Delhi. 1977.

7.     James E. Huheey, Inorganic Chemistry, Third Edition, Harper & Row, Publishers, Singapore 1983.

8.     Sanderson .R.T. Inorganic Chemistry, Affiliated East-West Press Pvt. Ltd, New Delhi. 1971.

9.     Seel F. Atomic Structure & Chemical Bonding, Methuen & Co., Ltd, 11 New Fetter Lane, London EC4. 1966.

10.  Wells A.F; Structural Inorganic Chemistry, Oxford University Press, Amen House, London EC4. 1962.

Problems:

1. Examine whether the following statements are true or false.

a)    In water sodium ion is more stable than sodium atom.

b)    Na²⁺ ion can be stabilized in a suitable medium.

c)     Sodium ion is more stable than sodium atom in vacuum.

d)    Ca⁴ ion can be stabilized in a suitable medium.

e)     There are stable O^2- and S^2- ions.

f)We cannot envisage a single NaCl molecule.

g)     The nucleophilicity of halide ions will never change with the change of solvant.

h)    There cannot be a single HCl molecule.

i) The binding energy of ions in a crystal is nearly equal to that in an isolated molecule.

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j)      CsCl does not adopt NaCl structure because Cs atom is highly electroposi­tive.

k)    In an ionic reaction,charge has greater influence than size.

l) In an ionic reaction,electronegativity difference has greater influence than size.

m)     As sodium is a reactive metal, one atom of sodium will readily react with an atom of chlorine to give two separate ions.

n)       In water medium, KBr will never react with AgF.

o)       AgF is insoluble in water.

p)       "Negative ions are less stable than positive ions in vacuum"

(q).   The stability of sodium ion in water is greater than the stability of sodium atom in water.

(r).     There can be no stable O^2- ion in solution or in a lattice.

2.     Will an atom of sodium react with an atom of chlorine if they are brought in contact with each other in vacuum'? Give reasons for your answer.

3.     Is it possible to stabilize Al⁺, Ca⁺ and Be⁺ ions?

4.     K²⁺ ion is not stable in water. Why?

5.     Why can there be no stable O^2- ion?

6.     Will the lattice energy stabilize O^2- ion?

7.     Is there solvation sheaths around anions in solution? Explain with an example.

8.     KBr does not adopt CsCl structure. Why?

9.     Though TiO₂ has rutile structure, description of the structure of TiCl₄ in terms of ionic bonding is not possible. Why?

10. Illustrate the statement "appropriate matching of the cation and the anion favours stability".

11. Justify the following reactions :

(a) 

Na₂SO₄ + BaCl₂                                     BaSO₄     + 2NaCI

(b)  K₂CrO₄ + Pb(OOCCH₃)₂                         2K(OOCCH₃) + PbCrO₄

12. Though interionic distances are almost the same, Na₂O (-602) has higher lattice energy than NaF (-214). Why?

13. Assuming that there can be no stable S^2- ion in solution, give the mechanism of precipitation of Hg²⁺,  Cu²⁺, Pb²⁺, Bi³⁺ and Cd²⁺ as sulphides in the IIA group.

14.    Though B³⁺ is isoelectronic with helium, it has never been observed in a compound. Why?

15.    The following passage is quoted from the "Chemical Bonding and the Geometry of Molecules" by George E. RYSCHKEW ITSCH. 1965 Page 46.

"As with simple positive ions, there is yet another limitation on the charge that can be carried out on a simple negative ion. When two negative charges are placed on an atom, the second electron must be added against the repulsion of the negative charge already on the atom. The result is that energy is absorbed in the process. The addition of a third electron to a doubly charged negative ion required a still greater amount of energy. As a consequence of this effect, the negative charge on simple ions is limited to a maximum of three units. For instance Si^4- ions are not found in compounds". Do you justify the above statement? Explain your answer.

16.    There is no limitation to the number of Na⁺ - Cl^- pairs that can combine to give solid NaCl. Why?

17.    Will an atom of sodium react with an atom of chlorine in vacuum to give two separate gaseous ions?

18.    Whereas Ba(NO₃)₂ is soluble in water, BaSO₄ is insoluble in water. Why? (Explanation on the basis of solubility product not expected).

19.    Examine which of the following statement is more reasonable. Il­lustrate your answer.

 (a) "The electron affinity of dinegative ions such as O^2- , S^2- is also negative. This means their enthalpy of formation is positive. Such ions cannot exist except through stabilization by environment, either in a crystal lattice or by solvation in solution"----

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From Inorganic Chemistry by James E. Huheey, Harper and Row. Pub­lishers, Inc, 100 East 53rd Street, New York, Ny 10022, 1983, Page 49.

(b) "Two and multi charge simple anions of O are not formed: compounds containing O^2- ions therefore do not exist in nature. Even in crystalline oxides of the Na₂O and CaO type, the effective charge of oxygen is only about 1 ". - From Inorganic Chemistry by N. AKHMETOV, MIR Publishers, Moscow.

20.   In water, HI is a stronger acid than HCl but  F ion is a better nucleophile

than      ion. Why?

21.   CaF is not a stable compound. Why?

22.   PbCl₂ and AgCl are insoluble in cold water. But PbF₂, AgF and BaCl₂ are soluble in water. Why?

23.   "For number of alkali halides, the radius ratio is lower than 0.414 (eg:LiI) but they adopt the NaCl Structure". Can you give an explanation for this.

24.   Reaction of the type

ICl + F₂                        IClF₂   is not probable. Why?

25.   Though BF₃ catalyses the ionization of R-F, it does not catalyse the ionization of R-Cl. Why?

26.   Which is more electronegative - fluorine atom or Na⁺ ion?

27.   The structure of the,ionic compounds does not depend on the nature or the electronic configuration of the element". Justify the above statement.

28.   "Whereas the packing in osmium metal is hexagonal or FCC in both of which each interior atom has 12 closest neighbours, the packing in tungsten metal is BCC in which each interior atom is attached directly to only eight others." Explain why does not tungsten adopt the hexagonal or F.C.C. packing as in osmium.

29.   Complex ions such as KCI^5- are not formed. Why?

30.   Most of the ionic crystals are insulators. Why?

This type of questions are suitable for open book system of examination.

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31.    The structure adopted by an ionic compound depends on the radii of its ions. Why?

32.    Though the ions in solution are more stable than the respective atoms, the reaction between ions are faster than the reaction between atom,. Why?

33.    Interhalogen compounds such as IBrCl₂, IBrF₂, IBrF₄ do not exist. Why?

34.    Some ionic solids do conduct electricity. What is it due to?

35.    Anhydrous CaSO₄ is slightly soluble in water but the dihydrate is a pure precipitate. Why?

36.    In a given medium several factors such as charge density, ionic size, polarizability, electronegativity difference, lattice, influence the course of the chemical reaction. With suitable examples give the order of influence of such factors on chemical reactions.